GCSE Combined Science Chemical energy changes revision

Exothermic reactions hand energy over to the surroundings, so the reaction mixture and everything touching it warms up: combustion, the neutralisation in chapter 15’s practical, and metals rusting or burning. Endothermic reactions soak energy up from the surroundings, which cools down: thermal decomposition needs continuous heating, and citric acid reacting with sodium hydrogencarbonate can chill a drink on a can’s base. The thermometer is the classifier — rises, exothermic; falls, endothermic.

Industry exploits both directions. Hand warmers ride exothermic chemistry: disposable ones let iron rust slowly in air (an oxidation), while rechargeable pouches hold supersaturated sodium ethanoate that releases heat the moment you click it into crystallisation — boiling the pouch afterwards redissolves the crystals to reset it. Self-heating cans mix calcium oxide (quicklime) with water; self-chilling cans run the citric-acid reaction on a compartment of the drink.

A reaction profile plots energy (y-axis) against the progress of the reaction. Reactants start on a ledge, energy climbs to a peak at the transition point, then drops to the products’ ledge. For an exothermic reaction the products’ ledge sits lower than the reactants’ — the spare energy left the reaction — and the ΔH arrow points down, drawn negative. For an endothermic reaction the products sit higher, and ΔH points up, positive.

(Higher) The height of the climb from the reactants’ ledge to the peak is the activation energy — the minimum energy colliding particles must carry for the reaction to happen. A catalyst offers an alternative route with a lower peak: it cuts the activation energy and speeds the reaction, but the two ledges do not move, so ΔH is unchanged.

(Higher) Every bond holds atoms together by storing energy, and the bond energy is what it costs to break one mole of that bond in the gas phase. Breaking bonds is therefore endothermic (energy in), and making bonds is exothermic (energy out). The overall energy change is the ledger’s bottom line: ΔH = energy to break all the bonds − energy released forming all the new bonds. Negative answer → exothermic; positive → endothermic.

The ledger explains chapter 15’s acids too: stronger bonds formed than broken means energy is left over for the surroundings, which is why neutralisation always warms the cup. It also explains why a catalyst cannot change ΔH — the ledger depends only on where you start and finish, not on which route (or lower peak) you travel.

This practical finds the exact reacting ratio of an acid and alkali by watching the temperature climb. Method: place 25 cm³ of dilute hydrochloric acid in a polystyrene cup (an insulator) with a lid, and record its starting temperature. Add 5 cm³ of sodium hydroxide solution, fit the lid, stir, and record the maximum temperature. Add a fresh 5 cm³ portion to the same mixture, stir, record again — and repeat, so the total volume of alkali grows by 5 cm³ each time.

Plot maximum temperature change against volume of alkali added. The graph shows two straight lines: a rising line while acid is in excess, and a falling line once alkali is in excess and simply adds cold volume. The intersection marks the volume of alkali that exactly reacts with all the acid — the maximum temperature change, because that is where the most neutralisation happens.

Hazards: dilute acid and alkali are irritants — goggles on, spills rinsed at once; hot or recently stirred mixtures can spit, so keep faces clear. A glass beaker loses far more heat, so results read low — the polystyrene cup and lid are accuracy equipment, not decoration.